Greetings. In this lecture I'll be discussing solubility rules for ionic compounds dissolving in water. First, I want you to recall the definitions of solubility. Solubility for ionic compounds was related to the amount of solid that would dissolve in the water. You can do this for other solvents as well. But, in the next few lectures we'll be considering aqueous solutions. We determine that if the concentration of ions can be above 0.1 molar then that species was defined as being soluble in that solvent. But without doing the experiments ourselves, without looking at all the different ionic compounds that are available and trying to see how well they dissolve in water, how can we predict whether or not a particular ionic compound will dissolve in water? Well, predicting something like this takes us all the way back to one of the first things we discussed in this course. And that is Coulomb's Law. Remember, Coulomb's Law dealt with the attraction or repulsion between charged particles. One of the first things I'd like to point out when we're discussing Coulomb's Law in the context of solutions is that for aqueous solutions we will have water as the solvent and the dielectric constant of water is relatively high. Because the dielectric constant is in the denominator and the dielectric constant of water at room temperature is about 80. [SOUND] We'll be dividing whatever is in the enumerator by a factor of 80 times the distance squared. That means that the force of attraction between ions in water is low relative to what that force of attraction would be in a lower dielectric medium. For example, air has a dielectric constant of about one, and even some other liquid solvents like some hexane or some carbon disulphide have relatively low dielectric constant. Carbon disulfide has a dielectric constant of only 2.6, which is a lot lower than 80. Now let's look at the numerator. If the charges, q1 and q2 are small in magnitude, in other words, if their absolute value is one, then we expect the force of attraction to also be relatively small. If we increase the magnitude of the charges, for example, if we go from plus one for an alkaline metal to plus two for an alkaline earth metal, or plus thre for aluminum, for example, then the force of attraction becomes stronger. Because the charges are on the numerator. Distance, on the other hand, is in the denominator. So, if you have a large distance between the particles, then the force of attraction is small. One of the ways that can happen with ionic compounds is to have an ion that has a relatively large atomic radius. Or to have a polyatomic ion, which has many atoms together, with, for example, a minus one charge spread out over a large distance. So let's keep Coulomb's Law in mind as we talk about the solubility rules. You can find various solubility rules listed in general chemistry textbooks or on webpages throughout the Internet. These rules are also summarized on the reference page of the course homepage. I'm going to list them here and then we'll do some examples. The first rule is that compounds of ammonium ion, or the 1A metal cations are soluble. Now, if you don't have periodic table handy you might not remember that the 1A metal cations include lithium, sodium, potassium, cesium, and rubidium. All of those have a plus one charge as their most frequent oxidation state. Exceptions to rule number one are fairly rare. So most of the time, ammonium compounds, sodium compounds, lithium compounds, potassium compounds, anything covered by rule number one have high solubility in water. Rule number two, is now looking at it from the perspective of the anion, whereas rule number one listed a bunch of cations, rule number two lists a bunch of anions. This rule is summarized as compounds of many atomed low charged polyatomic ions are soluble. What does that mean? Well here's some examples. If you have a whole bunch of non-metals grouped together, and that group of atoms has a negative one charge, and here's some examples, acetate, perchlorate, chlorate, and nitrate, then the ionic compounds of those species are soluble. So you could have, for example, aluminum acetate, or lead 2 nitrate, or copper 1 perchlorate, all of those compounds are soluble in water. It comes down to Coulomb's law once again. If we relate it to Coulomb's Law, the cation and anion attractive forces are relatively small when the charge values are small. And here we've got charge values of either plus one or minus one, depending on which rule we're looking at. We can summarize this as follows. Ionic compounds with ions of low charge are often soluble in water. Let's talk about a few more of these rules. Rule number three deals with halides. Again, halides have minus one charges. So chloride, bromide and iodide ion's compounds are soluble except those of silver 1, copper 1, thallium 1, mercury 1, which is often shown as this dimer, and lead 2-plus. Those are the very common exceptions. Rule number three doesn't quite contain all of the halides, does it? Fluoride salts, in contrast to what's listed in rule number three, are typically of low solubility in water. In fact, even lithium fluoride has a Ksp value of only 1.84 times 10 to the minus 3. That's a small Ksp value, right? It's less than 1, indicating that lithium fluoride is not particularly soluble in water. Why do you think this is the case? Rule number four says that compounds of sulphate ions are usually soluble. Notable exceptions include those compounds that contain calcium 2 plus, strontium 2 plus, barium 2 plus, or lead 2 plus. So, strontium sulfate is not soluble. But other compounds of sulfate are soluble. Sulfate now, has moved into the realm of having a minus two charge, but it has lots of atoms and it's relatively large. So the minus two charge is spread out over a large area. And so there's lots of compounds of sulfates that are soluble. The last rule is the catch all rule. And my last rule is that most other ionic compounds are insoluble. Sometimes a list of rules is longer. Because people have chosen to break up rule number five into more specific rules. For example, they might have a rule that says something like most silver salts are insoluble except and then give a bunch of exceptions. Or most hydroxide salts are insoluble except those of alkaline metals, etc. Most of the later rules on lists like this are simply redundant. So I prefer to try and keep the list at rule number 5 and deal with the few exceptions that occur from there. Now let's practice applying these solubility rules. Let's determine whether each of the following compounds is soluble in water. Why don't you go ahead and take out a piece of paper and write these down. You can look at the solubility rules. And you can look at the periodic table. And decide whether you think each of these is soluble or insoluble in water. You can pause the video now to write down your answers, and then we'll go through these one at a time. Let's start with the first item on the list, potassium hydroxide. Is potassium hydroxide soluble or insoluble in water? Well, if we look at the list of solubility rules, the first rule is that compounds of ammonium ions and the 1A metal cations are soluble. This of course would include potassium 1+ salts, so potassium hydroxide is soluble. How about the next compound on the list, silver bromide. Is silver bromide soluble? To answer this question about the solubility of silver bromide we have to go back to our list of solubility rules. There we find that rule number three talks about the relatively high solubility of most bromide compounds in water. Except it says explicitly that silver one bromide is not soluble. So most bromides are soluble except silver-1, copper-1, thalium-1, et cetera. So silver bromide is an exception, and it is insoluble. That's how we use those exceptions when we're looking at the rules. Let's move to the next compound on our list of examples. Calcium chloride. Is calcium chloride soluble in water? Again, it's rule number three that applies from our list of solubility rules. Most chloride compounds are soluble. And I don't see calcium listed as an exception. So I can assume safely then that calcium chloride is soluble in water. How about the next example, lead 2 nitrate. Is lead 2 nitrate soluble in water? This example uses solubility rule number two, and that is that compounds of these many atom, low charge polyatomic atoms are soluble. And that includes nitrates. There aren't any exceptions listed so we can safely assume that lead 2 nitrate is soluble. Finally, what about lead 2 sulfate? Which of our rules applies to this compound? Is lead 2 sulfate soluble in water? There is a rule that applies to the solubility of sulfate specifically and it says that most sulfate compounds are soluble. But lead 2 sulfate is listed as an exception so lead 2 sulfate is insoluble. If you'd like to practice more of these problems, you can think of any ionic compound you can think of and look at the list of solubility rules and see which of the rules apply. And then, you could even Google the Ksp values and find out what the actual solubility product constant is of that compound if it turns out that it's insoluble. For example, I looked up the solubility product constant of silver 1 bromide and found that it was a very, very small number, which means that silver 1 bromide is extremely, sparingly soluble in water. Very, very few ions will dissolve, even if you have a very large amount of water. I also looked up the Ksp value for lead 2 Sulfate. And again, it's a small number. Much, much less than one. This concludes the lecture on applying solubility rules. We're going to use these rules as guidelines in the next few lectures, as we predict whether or not a precipitate will form when we mix aqueous solutions.
